Question

For the spontaneous reaction $\Delta G$ , equilibrium K and E cell will be respectively:
A.$- ve > 1, + ve$
B.$+ ve, > 1 - ve$
C.$- ve, < 1, - ve$
D.$- ve, > 1,ve$

Answer 1

The sign of the standard free energy change ‘ G ‘ of a chemical reaction determines whether the reaction will proceed in the forward or reverse direction.

Complete step by step answer:
1.We know that $\Delta G$ is only meaningful for changes in which temperature and pressure are constant. These are the conditions under which most reactions occur in the laboratory; the system is usually open to the atmosphere (constant pressure).
2.We begin and finish the process at room temperature (after any heat that we have added or that has been released from the reaction).
3.The significance of the Gibbs function can hardly be overstated: it acts as a single master variable that determines whether a given chemical change is thermodynamically possible.
4.Thus if the free energy of the reactants is greater than that of the products, the entropy of the world will increase as the reaction proceeds, and therefore the reaction will spontaneously occur.
5.Conversely, if the free energy of the products exceeds the reagents, the reaction will not take place in the written direction, but it will proceed in the reverse direction.

Hence, the correct option is D.

Note:
In a spontaneous change, the Gibbs energy always decreases and never increases.

An important consequence of the one-way path of free energy is that once it reaches its lowest possible value, all net changes stop. This, of course, represents a state of chemical equilibrium. These relationships are as follows:

To $G < 0$: Response may proceed spontaneously to true:
$A \to B$
$\Delta G > 0$: The reaction may spontaneously proceed to the left:

$\Delta G = 0$: reaction is at equilibrium; The quantity of [A] and [B] will not change
Remember the situation for easy change
$\Delta G = \Delta H - T\Delta S < 0$.